Chemistry Grade 11 Study Notes

  • CHEMISTRY STUDENT TEXTBOOK GRADE 11: This is a chemistry textbook for grade 11 students.\
  • Authors and Editors: Lists the writers, editors, illustrator, and designer involved in creating the textbook.\
  • Copyright Information: Specifies the copyright year, the owner (Federal Democratic Republic of Ethiopia, Ministry of Education), and the rights reserved. It mentions the prohibition of reproducing or transmitting the textbook without permission, as per the Federal Negarit Gazeta, Proclamation No. 410/2004.
  • Acknowledgements: Expresses gratitude to individuals, groups, and bodies involved in publishing the textbook. Special thanks are given to Hawassa University and other universities like Addis Ababa University, Bahir Dar University, and Jimma University for their contributions.\
  • Contract and ISBN Information: Includes printing details, contract number (MOE/GEQIP-E/LICB/G-01/23), and the ISBN (978-99990-0-018-5).\
  • Book Care Instructions: Provides ten guidelines on how to protect the textbook from damage, such as covering it with protective material, keeping it in a clean place, and avoiding writing or tearing pages.\
  • Atomic Theory Development:
    • Ancient Greek philosophers (e.g., Democritus) pondered whether matter was continuously divisible or had an ultimate limit.
    • Democritus suggested matter was composed of tiny, indestructible particles called "atomos" (atoms), meaning "indivisible".
    • Dalton's atomic theory (1808) marked a significant step, gaining broad acceptance.
  • Scientific Laws:
    • Scientific laws often emerge from prior scientific findings.
    • Dalton's atomic theory is grounded in the law of conservation of mass and the law of definite proportions.
  • Example 1.1:
    • Water is always 11.2% hydrogen and 88.8% oxygen by mass.
    • a. In 18.0 g of water: mass of hydrogen = 0.112×18.0g=2.02g0.112 \times 18.0 g = 2.02 g, mass of oxygen = 0.888×18.0g=15.98g0.888 \times 18.0 g = 15.98 g
    • b. In 1.00 g of water: mass of hydrogen = 0.112×1.00g=0.112g0.112 \times 1.00 g = 0.112 g, mass of oxygen = 0.888×1.0g=0.888g0.888 \times 1.0 g = 0.888 g
  • Law of Multiple Proportions:
    • Example data for nitrogen and oxygen compounds (A, B, C) illustrate this law.
    • Ratios of nitrogen masses combining with 1 g of oxygen should be small whole numbers.
    • Ratios are computed as A/C = 4, B/C = 2, ensuring the Law of Multiple Proportions holds.
  • Dalton's Atomic Theory Limitations:
    • Some of Dalton's postulates are inconsistent with later findings.
    • Dalton's model remains useful despite inconsistencies, focusing mainly on how atoms combine.
  • Modern Atomic Theory Development:
    • Began with J.J. Thomson's discovery of the electron in 1897.
    • Modern atomic theory focuses on the physical structure of atoms, evolving through experiments and observations.
  • Cathode Rays:
    • J.J. Thomson's experiment (1897) involved high-voltage source sealed in a glass tube.
    • Cathode: Negative electrode
    • Anode: Positive electrode
    • Cathode rays originate from the cathode and interact with the glass, emitting a greenish light.
  • Properties of Electrons:
    • Thomson concluded that cathode rays consist of negatively charged particles (electrons).
    • Electron characteristics are independent of the cathode material.
    • Thomson calculated the mass-to-charge ratio of an electron: −5.686×10−12kg⋅C−1-5.686 \times 10^{-12} kg \cdot C^{-1}.
    • Millikan's oil drop experiment (1909) determined the charge of an electron: e=−1.602×10−19Ce = -1.602 \times 10^{-19} C.
    • Electron's mass was calculated to be 9.109×10−31kg9.109 \times 10^{-31} kg.
  • Radioactivity:
    • Radioactivity (radioactive decay): spontaneous emission of particles and/or radiation from unstable nuclei.
    • Three types of rays identified:
      • Alpha ($\alpha$) rays: Positively charged particles, identical to helium nuclei.
      • Beta ($\beta$) rays: Electrons coming from inside the nucleus.
      • Gamma ($\gamma$) rays: High-energy rays with no charge, unaffected by external fields.
  • Discovery of the Nucleus:
    • Thomson's plum-pudding model: electrons and protons randomly distributed in a positively charged cloud.
    • Rutherford's experiment: positively charged particles aimed at a thin sheet of gold foil.
    • Results: Some particles deflected, a few deflected backward.
    • Conclusion: Atom consists mostly of empty space with a small, dense, positively charged nucleus.
      Neutron:
    • British physicist James Chadwick (1891–1974) demonstrated that beryllium metal, when irradiated with alpha rays, emits strongly penetrating radiation, and showed that this radiation consists of neutral particles, called neutrons.
  • Make-up of the Nucleus:
    • (1919) Rutherford discovered protons (hydrogen nuclei) form when alpha particles strike nitrogen.
    • Proton: positively charged nuclear particle with charge equal to electron's magnitude; mass mp=1.67262×10−27kgm_p = 1.67262 \times 10^{-27} kg.
    • Nucleus's positive charge is due to protons.
  • Subatomic Particles:
    • Properties table:
      • Particle, Actual mass (kg), Relative mass (amu), Actual charge (C ), Relative charge
      • Proton (p), 1.672622×10−271.672622 \times 10^{-27}, 1.007276, 1.602×10−191.602 \times 10^{-19}, +1
      • Neutron (n) ,1.674927×10−271.674927 \times 10^{-27}, 1.008665, 0, 0
      • Electron (e−) ,9.109383×10−319.109383 \times 10^{-31}, 5.485799×10−45.485799 \times 10^{-4}, -1.602×0−191.602 \times 0^{-19}, -1
    • Atomic number (Z): # of protons in nucleus; defines the element.
    • Mass number (A): total number of protons and neutrons in nucleus.
    • Notation:
  • Isotopes:
    • Atoms of the same element with different neutron numbers and mass numbers.
    • Natural elements are often isotopic mixtures.
    • Average atomic mass: weighted average of isotope masses.
    • Formula: A=A<em>1f</em>1+A<em>2f</em>2+…+A<em>nf</em>nA = A<em>1 f</em>1 + A<em>2 f</em>2 + \ldots + A<em>n f</em>n, where Ai are the relative masses and fi are fractional abundances.
      Electromagnetic Radiation:
    • James Clerk Maxwell (1873) proposed light consists of electromagnetic waves.
    • Electromagnetic wave: electric field and magnetic field components vibrating in perpendicular planes.
    • EMR: emission and transmission of energy in electromagnetic waves.
    • Characteristics: wavelength ($\lambda$), frequency ($\nu$), and speed (c).
      Electromagnetic Spectrum:
    • Wavelength ($\lambda$): Distance the wave travels during one cycle, expressed in meters (m) or nanometers (nm).
    • Frequency ($\nu$): Number of cycles per second, expressed in hertz (Hz) or 1/s1/s.
    • Speed (c): Product of wavelength and frequency: c=νλc = \nu \lambda.
    • In a vacuum, all electromagnetic waves travel at the same speed, which is the speed of light,(3×108m/s)(3 \times 10^8 m/s).
    • EMR frequencies span from radio waves to gamma rays.
      Quantum Theory and Photon:
    • Max Planck (1900) proposed that energy is discontinuous and emitted or absorbed in discrete quantities called "quanta".
    • Energy E of a single quantum given by: E=hνE = h \nu, (hh: Planck's constant 6.63×10−34J⋅s6.63 \times 10^{-34} J \cdot s).
    • Since ν=c/λ\nu = c/\lambda, E=hc/λE = h c/\lambda .Thus, energy has particulate properties.
      Photoelectric Effect:
    • Albert Einstein (1905) explained photoelectric effect using quantum theory.
    • Photoelectric effect: Emission of electrons from the surface of certain metals exposed to light of at least a specific minimum frequency (threshold frequency).
    • Minimum energy to remove electron: E<em>0=hν</em>0E<em>0 = h \nu</em>0.
    • Kinetic energy of ejected electron: KE<em>e=1/2mv2=hν−hν</em>0KE<em>e = 1/2 m v^2 = h \nu - h \nu</em>0.
    • Energy is Mass: E=mc^2 and apparent mass of Photon=hcλ\frac{h}{c \lambda}
      Bohr Model of Hydrogen atom
    • Bohr proposed that the absorptions and emissions in line spectra correspond to the transfer of the electron from one orbit to another.
    • An electron in hydrogen atom travels around the nucleus in a circular orbit.
    • The energy of the electron in an atom is proportional to its distance from the nucleus. The further an electron is from the nucleus, the more energy it has.
    • Only limited number of orbits with certain energies are allowed. This means, the orbits are quantized.
    • Radii, r, of permitted orbits for hydrogen related to Planck’s constant, h, electron’s charge, e, and mass, m.
    • E=mv2−<br/>Calculationofradius:</li><li>Permittedorbitradiiproportionaltothesquareofawholenumber,n:E= mv^2 -<br /> Calculation of radius:</li> <li>Permitted orbit radii proportional to the square of a whole number, n:r = n ao,,, ao=0.53 A</li><li>Energiesthatanelectroninhydrogenatomcanoccupyaregivenby:</li><li></li> <li>Energies that an electron in hydrogen atom can occupy are given by:</li> <li> En = \frac{-RH}{n^2} ,,RH=2.18 X10^{-18}J.<br/>Balmer/LymanSeriesinhydrogenatomicspectra.<br/>DualityofmatterandEnergy:</li><li>DeBrogliederivedanequationforthewavelengthofanyparticleofmass,m,whetheraplanet,ball,orelectron−movingatspeed,v:.<br /> Balmer/Lyman Series in hydrogen atomic spectra.<br /> Duality of matter and Energy:</li> <li>De Broglie derived an equation for the wavelength of any particle of mass, m, whether a planet, ball, or electron-moving at speed, v:\lambda = \frac{h}{mv} <br/>HeisenberguncertaintyPrinciple</li><li>Itisnotpossibletoknowwithgreatcertaintybothanelectron’spositionanditsmomentump(wherep=mν)atthesametime:(<br /> Heisenberg uncertainty Principle</li> <li>It is not possible to know with great certainty both an electron’s position and its momentum p (where p = mν) at the same time: (\Delta x )()(\Delta$\p)≥ h4π\frac{h}{4\pi}
      Quantum Numbers:
    • Erwin Schrödinger (1927), suggested that an electron or any other particle exhibiting wavelike properties can be described by a mathematical equation called a wave function (denoted Greek latter psi, ψ).
      -Principal Quantum Number (n):It may be any positive integer, n = 1, 2, 3, 4, etc. It describes the size and energy of the shell in which the orbital resides and it is analogous to the energy levels in Bohr’s model.
      -The angular momentum quantum number (ℓ) designates the shape of the atomic orbitals. It takes values from 0 to n-1. Orbitals of the same n, but different ℓ are said to belong to different subshells.
      The magnetic quantum number (mℓ) is also called the orbital-orientation quantum number. It has integral values between -ℓ and ℓ, including 0.
      -The electron spin quantum number (ms) refers to the spin of an electron and the orientation of the magnetic field produced by this spin. It takes the value +½ or -½
      Shapes of Atomic Orbitals:
    • Three-dimensional aspects of atomic orbital orientation are represented by boundary surface diagrams.
    • The region of greatest probability for finding an electron in an s orbital is spherically symmetrical.
    • The p orbitals are arranged along three mutually perpendicular axes (px, py, pz), with dumbbell-shaped boundary surfaces.
    • d Orbitals: More complex shapes and spatial orientations compared to p orbitals. n ≥ 4 have additional fourth sublevels, the f orbitals have more complex shapes than the d orbitals
      Aufbau priciple:This is a scheme used to reproduce the electron configurations of the ground state of atoms by successively filling with electrons in a specific order (the building up order). In general, electrons occupy the lowest-energy orbital available before entering the higher energy orbital. Accordingly, the ground state electron configurations of atoms are obtained by filling the subshell in the following order; 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p,5s, 4d, 5p, 6s, 4f, and so on
      Hund’s Principle: Equal energy orbitals (degenerate orbitals) are each occupied by a single electron before the second electrons of opposite spin enters the orbital. In other words, each of the three 2p orbitals (2px, 2py and 2pz) will hold a single electron before any of them receives a second electron.
      Pauli’s Exclusion Principle: No two electrons can have the same four quantum numbers. That means, they must differ in at least one of the four quantum numbers.
      Periodic Law: Certain sets of physical and chemical properties repeat at regular intervals when elements are arranged by increasing atomic number.
      Trends such as :Atomic Size (Atomic Radii), Ionization Energy (IE), Electron Affinity (EA)Electronegativity and Metallic Character
      Chemical Bonds:
    • Ionic bond: electrostatic attraction between positive and negative ions.
    • Covalent bond: formed by the sharing of a pair of electrons between two atoms.
    • Metallic bond: the sharing of free electrons (delocalized electrons) among a lattice of positively charged metal ions.
      Properties of ionic Compounds
      -Ionic compounds are usually formed when metal cations bond with non-metal anions.
      -lattice energy (U) is the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions LATTICE ENERGIES FROM THE BORN–HABER CYCLE are used to cal. values by use of Hess Law.
      -Born–Haber cycle, useful when working with lattice energies. The reasoning is based on Hess’s law, which states that an overall reaction’s enthalpy change is the sum of the enthalpy changes for individual reactions: ∆H<em>total=∆H</em>1+∆H<em>2+∆H</em>3+…∆H<em>\text{total} = ∆H</em>1 + ∆H<em>2 + ∆H</em>3 + …
  • The formation of ionic bonding is influenced by Ionization energy (IE),Electron affinity (EA)and Lattice energy Lewis Structure:
    • Consists of a chemical symbol, representing the nucleus and core electrons, surrounded by dots for valence electrons.
    • Lewis structures illustrate electron transfer or sharing in chemical bonds.
      Coordination Covalent Bonds and Resonance Structures:
    • Coordinate covalent bonds involve one atom donating both electrons to another atom that has a vacant valence orbital
      Electronegativity of elements: Electronegativity generally increases across a period from left to right (say from lithium to fluorine) and decreases down a group
      Metallic Character: Metallic character decreases as you move across a period in the periodic table from left to right. Metallic character increases as you move down an element group in the periodic table.
      Exception to Octet Rule: Atoms of H, Li, Be and B attain an arrangement of two electrons like He.,the transition and post-transition elements do not usually obey the octet rule. For For transition metals, the 18-electron rule replaces the octet rule . In General, there is a gradual change from metallic to non-metallic character as you move from left to right across a period and from bottom to top within most groups in the periodic table. Accordingly, atoms can combine to form three types of bond: metal with non-metal (ionic bond), non-metal with non-metal (covalent bond), and metal with metal (metallic bond).
      Electrical conductivity: to test the electrical conductivity of the aqueous solutions of some common ionic compounds. Apparatus and chemicals: 9-volt battery, 6-watt bulb with a bulb holder, conducting wires, two carbon rods, H2O, lead (II) iodide, NaCl. Safety lead compounds are harmful. Covalent Bond:
      -Formed when a pair of electrons is shared between two atoms due to: attraction and Repulsion Forces
      -Lewis Structures:Lewis structures depict covalent bonding showing bonding and non-bonding pairs.The steps to writing a Lewis structure can be presented in a flow diagram
      -Coordinate-Covalent Bonding:Coordinate Bond Formation and Coordinate Compounds: coordinate bond formation from ammonia and boron trifluoride with the octet rule shown.
      Exceptions to Octet Rule: odd Numbers of electrons in molecules: there are are some examples like ClO2, NO and NO2 where there are bond the atoms in an the molecule violates octer rule
      Polar and non polar covalent molecules-electronegativity: Difference in electronegativity results in unequal sharing of electrons, and bond polarity and dipole
      Dipole Moment: Diatomic Molecules, Polyatomic Molecules
      Metallic bond:Metallic bonds are the chemical bonds that hold atoms together in solid metals
      Based on the type of subshell being filled elements are divided into 3 group of elements
      Several Questions that raised while looking at the electron configuration of an atom are answered by the followings principles:
      Aupbau principle, Hund Principle,Pauli exclusion principle